Determining the enthalpy of a reaction

Determine the enthalpy of this reaction:

Mg(OH)2(s) + H2SO4(aq) → MgSO4(aq) + 2H2O(l)

Method

Measure out 50 mL of 1 mol L-1 sulfuric acid (an excess) into a styrofoam cup.

Weigh accurately about 1.5 g of Mg(OH)2 powder.

Measure the temperature of the sulfuric acid.

Quickly add the Mg(OH)2 powder, stir and measure the maximum temperature of the reaction mix.

Measure 50 mL of about 1 mol L-1 sulfuric acid.

Pour the acid into a styrofoam cup (two stacked together will give even better insulation).

Weigh accurately about 1.5 g of Mg(OH)2 powder.

Measure the initial temperature of the sulfuric acid solution.

Tinitial = 16.5 °C

Add the Mg(OH)2 to the acid.

Stir well and watch the temperature. All the powder should dissolve since we have excess acid.

Measure the maximum temperature change.

Tfinal = 28.0 °C

Calculate:

• The amount, in moles, of Mg(OH)2 reacting.

• The total energy released during the reaction.

• The energy change per mole of Mg(OH)2 reacting.

The energy released by the reaction has made the temperature of the solution rise by

28.0 °C – 16.5 °C = 11.5 °C

All dilute aqueous solutions — and 1 mol L-1 is still considered to be dilute — contain very much more water than they do any other reagent. We therefore assume that the specific heat of the solution is equal to that of water, which is 4.18 J °C-1 g-1. That is, it takes 4.18 J of energy to change the temperature of 1 g of the solution by 1 °C.

The mass of 1 mL of water at room temperature is 1 g. We assume that 50 mL of acid has a specific heat equivalent to 50 g of water.

Energy =

mass of

water heated

×temperature change

× 4.18 J °C-1 g-1

Mg(OH)2(s) + H2SO4(aq) → MgSO4(aq) + 2H2O(l)

∆H = -92.9 kJ mol-1

The temperature rose, so the reaction is exothermic and the ∆H is negative: